Chemistry Chapter 7 - summary PDF

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Chemical bonds are formed because they reduce the potential energy between the charged particles that compose atoms. In essence, the Potential energy of bonded atoms < Potential energy of separate atoms. E.g: In the bond formation of Hydrogen (H2), the potential energy of the two separate hydrogen atoms reduces as they get closer to each other to form a covalent bond. The atoms cannot get too close to each other, else the energy will increase and make it unstable (according to Coulomb’s law). Thus the bond length is the internuclear distance at which the lowest potential energy is attained.

Ionic bonds: ● Formed by the electrostatic force between the cation(+) and anion(-) of the metal and non-metal respectively. ● Metals have low ionization enthalpies and form cations easily by losing electrons ● Non-metals have high electron affinities and form anions easily by losing electrons. ● Cations attract anions to lower overall potential energy. ● Ionic solids are crystalline, rigid, brittle and strong. They have high melting and boiling points - because of the strong attraction forces caused by ionic bonding. In the solid state, they are poor conductors of electricity. But when they are in their melted or dissolved state, they are excellent conductors of electricity. They readily dissolve in water. Covalent bonds: ● Formed by the sharing of electrons between two nonmetals ● Non-metals have high ionization enthalpies and high electron affinities. Thus they have to share. ● The shared electrons interact with the nucleus of both atoms and thus lowers the potential energy and hence makes it more stable. ● Covalent compounds have low melting and boiling points. They exist as liquids and gases at room temperature. They are insoluble in water and are poor conductors of electricity. Some covalent solids are softer than ionic solids. (exception diamond)

Ionization Enthalpy/Energy - Energy required to remove an electron from its last shell. Electron Affinity - amount of energy released or spent when an electron is added to a neutral atom or molecule in the gaseous state to form a negative ion. Polar Covalent Bonding:

● Involves assigning a partial positive and a partial negative charge to the extreme ends.(Metals get the positive and nonmetals get the negative) ● Electron density is not the same on both the atoms. ● This type of bonding is an intermediate between ionic and pure covalent bonding. ● The partial negative charge indicates the electron density at that specific atom.

Non-polar Covalent Bonding: ● There is equal distribution of electron on both sides of the compound. (Electron density is the same) ● A common example: Hydrocarbons (oil) Metallic Bonding: ● Metals bonding to metal free electrons circulating between them, thus making them conducive. ● The absence of rigid ionic bonds make it malleable.

Electronegativity: It is the tendency of the atom to attract bonding electrons in a chemical bond. It increases across a period and decreases down the group. Fluorine is the most electronegative element. Francium is the least electronegative element. It is not the same as electron affinity. In an electron distribution map, the non-metals have a higher electron density due to the property of electronegativity. These regions are marked in red and the regions with lesser electron density( less electronegative atoms) have their areas marked in blue. Polarity ∝ Electronegativity difference. I.e Greater the electronegativity difference means a greater polarity. Another measurement of polarity is dipole moment. Octet rule: Tendency of main group elements to form enough bonds to obtain 8 electrons in its outermost shell. Single bond: When two elements share one pair of electrons Double bond: When two elements share two pairs of electrons Triple bond: When two elements share three pair of electrons. Length trend of bonds: Triple< Double< Single Strength of bonds: Triple> Double< Single

Drawing Lewis Structures:

1. Identify which atom should be in the centre of the structure. The least electronegative atom should always be there in the centre. Hydrogen atoms must always be kept in the periphery. 2. Calculate the total number of valence electrons in the compound. 3. Construct a single bond between the central atom and the rest of the atoms. 4. Distribute the rest of the electrons among the rest of the atoms giving octets or duets(to hydrogen) 5. If an octet is not formed, trying adding double and triple bonds. *The triple bond in N2 make Nitrogen unreactive. Lewis’ theory of covalent bonding: 1. It explains why some combinations of atoms don’t form molecules - because they need to form octets or duets. 2. It implies that atoms between atoms are directional. 3. The melting and boiling points of molecular compounds are relatively low - they have weak intermolecular forces. 4. Covalent compounds are poor conductors of electricity as there are no charged particles. 5. Lewis theory predicts that atoms are really stable if they form an octet of valence electrons. Some atoms violate that rule (e.g: Hydrogen - 2;Be - 4; B- 6) Exceptions to the octet rule: 1. Incomplete octets: Molecules or ions with less than 8 electrons around an atom. E.g: B, Be, Al 2. Expanded octets: Molecules or ions with less than 8 electrons around an atom. E.g: 3-period and beyond due to d-orbitals. 3. Odd-electron species: Molecules or ions with an odd number of electrons. They are called free radicals. Resonance: A phenomenon in which two or more Lewis structures can be drawn for the same compound.

Formal Charge: A system that allows us to distinguish between the alternating Lewis structures. The smallest seperation of the charges often gives us the correct shape of the structure. ,

Formal Charge: (no. of valence electrons)-(no. Of non-bonding electrons)-((½)*bonding electrons).. The sum of formal charges of all atoms in an ion must be equal to the charge of the ion. Resonance hybrid: Average of resonance forms shown by a Lewis structure. Resonance structures do not exist in nature, but a resonance hybrid exists in nature. There can be more than one Lewis structure that exists in nature. These three bonds are equivalent. Each bond is intermediate in strength and length between a double bond and a single bond. Bond order: Equals the ratio of the number of chemical bonds to the number of groups attached to the central atom.

Lattice Energy: Energy involved when seperated gas ions are packed to form an ionic solid. It is exothermic because of the attractive force between cations and anions. Hess Law: The change in the overall enthalpy of a stepwise process is the sum of the enthalpy changes of the steps. The most common application of Hess Law is Born-Haber Cycle. It breaks the formation of an ionic solid into various steps. 1. Sublimation energy - The heat required to sublime one mole of the substance at a given combination of temperature and pressure, usually at STP. 2. Ionization energy - The energy required to remove the most loosely bound electron from a gaseous atom. 3. Bond energy - The energy to break one mole of the bond in gaseous phase. 4. Electron affinity - The energy associated with the gaining of an electron by an atom in the gaseous state. 5. Lattice energy - The energy change for the process of adding an electron to a gaseous atom to form an anion. 6. Standard enthalpy of formation (total) - An enthalpy change for a reaction in which a mole of a pure substance is formed from its constituent elements at STP. 1+2+3+5= formation enthalpy Trends in Lattice Energy: 1. The lattice energy increases as we go down the group. 2. The lattice energy is inversely proportional to the stability of the ionic compound (the more negative the lattice energy is the more stable it is. ) This is because as the ionic radii increases down the group, the potential energy increases and this makes the ionic compound more stable. 3. Lattice energy decreases with the increasing magnitude of the ionic charge. This means that more heat is released. 4. Greater charged ions and smaller sized ions lead to more negative lattice energy levels (lesser).

Bond Energies: The bond energy of a chemical bond is the energy required to break 1 mole of a covalent bond in the gaseous phase. It is always positive and is specific to the bonding atoms. Trends in Bond Energies: 1. Bond energy and bond strength increases as the number of bonds increase. Eg: C≡C (837 kJ)> C=C(611 kJ)> C-C(347 kJ) C≡N (891 kJ)> C=N(615 kJ)> C-N(305 kJ) 2. Bond energy and bond strength increases as the length of the bond decreases. - Bonds get weaker as we go down a group due to the increase in radius and hence weaker attraction. - Bonds get stronger across a period due to decrease in radius and stronger attraction and increased electronegativity. Bond Order: Number of chemical bonds between a pair of atoms. As bond order increases, the average bond length decreases and the average bond energy increases.

Bond length is the distance between the nuclei of two bonded atoms. Average bond length represents the average length of a bond between two particular atoms in a large number of compounds. Trends in Bond Lengths: 1. The more electrons two atoms share, the shorter the covalent bond. E.g: c(triple bond)-o...